Assigning Oxidation Numbers (College Board AP® Chemistry): Revision Note
Assigning oxidation numbers
Oxidation numbers are a tool chemists use to track the electrons gained or lost during a redox reaction
Each atom in a neutral compound or ion is assigned an oxidation number
Oxidation numbers are the “charges” that atoms will have, if the shared electrons in a bond were assigned to the atom with the greatest attraction for those electrons
Rules for assigning oxidation numbers
The oxidation number of an atom that is a free element is 0
E.g. Copper (Cu) has 0
Diatomic elements also have an oxidation number of 0
E.g. Chlorine gas (Cl2) has 0
In ions of just one atom, the oxidation number is equal to its charge
E.g. Ca2+ has +2 and Cl- has −1
The sum of the oxidation numbers of all the atoms in a neutral compound is 0
E.g. H2O has 2 hydrogen atoms with +1 and 1 oxygen atom with −2
2(H oxidation number) + 1(O oxidation number) = 0
2(+1) + 1(−2) = 0
0 = 0
The sum of the oxidation numbers of all the atoms in a polyatomic ion is equal to the charge of the ion
E.g. NO3- has 1 nitrogen atom with +5 and 3 oxygen atoms with −2
1(N oxidation number) + 3(O oxidation number) = charge of the ion
1(+5) + 3(−2) = −1
−1 = −1
Metals always have positive oxidation numbers in compounds
Group 1 metals always have +1
Group 2 metals always have +2
Nonmetals can have positive or negative oxidation numbers, depending on the compound they are in
In most scenarios, hydrogen tends to have +1 oxidation number
The negative oxidation number of nonmetals from Group 15 to 17 can be calculated using the formula: negative oxidation number = valence electrons − 8
The table below shows common oxidation numbers for nonmetals
Table showing the negative oxidation numbers for nonmetals from Group 15 to 17
Nonmetal | Valence electrons | Oxidation number |
|---|---|---|
Group 17 | 7 | −1 |
Group 16 | 6 | −2 |
Group 15 | 5 | −3 |
The oxidation number of any element can be figured out using rules 4 and 5, once all other oxidation numbers have been assigned
In most formulas, the element with the positive oxidation number is written first (e.g. NaCl, CO2). This is only a guide: in OF2, oxygen is +2
Comparing oxidation numbers before and after a reaction identifies the species oxidized and the species reduced
If an atom’s oxidation number increases, its species is oxidized, e.g. in Mg + 2H+ → Mg2+ + H2, Mg goes from 0 to +2, so Mg is oxidized
If an atom's oxidation number decreases, its species is reduced, e.g. H goes from +1 to 0, so H+ is reduced
Examiner Tips and Tricks
Write the sign of an oxidation number before the number (+2) and the charge of an ion after it (Ca2+). For a one-atom ion the two values are equal, but in a polyatomic ion they are not: Mn in MnO4- is +7, not −1
An oxidation number belongs to one atom: each O in SO42- is −2, not −8
Read the formula before you assign. An uncombined element, e.g. Ag (s), is always 0. A halogen is −1 only when it is combined with a less electronegative element: in NaOCl (hypochlorite, OCl-), Cl is +1, not −1
Worked Example
Assign the oxidation number of all the atoms in the following compounds or ions:
CO2
SO42-
Answer 1:
Oxygen is −2 in most compounds (rule 7, Group 16; the exceptions are peroxides and OF2)
Therefore O oxidation number = −2
Using rule 4, the sum of all the oxidation numbers of the elements in a neutral molecule must be 0
1(C oxidation number) + 2(O oxidation number) = 0
1(C oxidation number) + 2(−2) = 0
1(C oxidation number) − 4 = 0
1(C oxidation number) = 4
C oxidation number = +4
Answer 2:
Oxygen is −2 in most compounds (rule 7, Group 16; the exceptions are peroxides and OF2)
Therefore O oxidation number = −2
Using rule 4, the sum of all the oxidation numbers of the elements in a polyatomic ion must be equal to its charge
Sulfate has a charge of 2−, so the sum must be equal to −2
1(S oxidation number) + 4(O oxidation number) = −2
1(S oxidation number) + 4(−2) = −2
1(S oxidation number) − 8 = −2
1(S oxidation number) = −2 + 8
S oxidation number = +6
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